Transformation of Reactive IronMinerals in a Permeable ReactiveBarrier (Biowall) Used to Treat TCEin GroundwaterY . T H O M A S H E , * , † J O H N T . W I L S O N , A N DR I C H A R D T . W I L K I N

U.S. Environmental Protection Agency, National RiskManagement Research Laboratory, Ground Water andEcosystems Restoration Division, 919 Kerr Research Drive,Ada, Oklahoma 74820

Received April 25, 2008. Revised manuscript received June18, 2008. Accepted June 18, 2008.

Iron and sulfur reducing conditions generally develop inpermeable reactive barrier systems (PRB) constructed to treatcontaminated groundwater. These conditions allow formationof FeS mineral phases. FeS readily degrades TCE, but atransformation of FeS to FeS2 could dramatically slow the rateof TCE degradation in the PRB. This study uses acid volatilesulfide (AVS) and chromium reducible sulfur (CRS) as probes forFeS and FeS2 to investigate iron sulfide formation andtransformation in a column study and PRB field study dealingwith TCE degradation. Solid phase iron speciation showsthat most of the iron is reduced and sulfur partitioningmeasurements show that AVS and CRS coexist in all samples,with the conversion of AVS to CRS being most significant inlocations with potential oxidants available. In the column study,54% of FeS was transformed to FeS2 after 2.4 years. In thefield scale PRB, 43% was transformed after 5.2 years. Microscopyreveals FeS, Fe3S4 and FeS2 formation in the column system;however, only pyrite formation was confirmed by X-ray diffraction.The polysulfide pathway is most likely the primary mechanismof FeS transformation in the system, with S0 as an intermediatespecies formed through H2S oxidation.

IntroductionPrevious studies on abiotic reduction of chlorinated solventshave examined different iron minerals, including Fe(0) (1),Fe(II) sorbed to iron oxides (2), green rust (3), magnetite (4),and iron sulfides (5, 6), and almost all of these studies areconducted in laboratory batch studies with synthesizedmaterials. There are a few field related studies providingevidence that iron sulfides have an important contributionto the degradation of chlorinated solvents in natural andengineered systems (7, 8).

Permeable reactive barrier (PRB) technology is commonlyused to treat groundwater contamination for both organicand inorganic contaminants (9–13). Both iron and sulfateare abundant in groundwater environments. Under the sulfurand iron reducing conditions typically encountered in PRBsystems, the formation of various FeS phases would beexpected, and FeS precipitation has been reported in several

reactive barrier systems (10, 12, 13). Amorphous FeS andpoorly crystalline mackinawite (tetragonal structure) havebeen observed in PRB systems constructed with zerovalentiron (ZVI) and in biowall systems constructed with plantmulch (12, 13). However, iron mineralogy in these systemsis not well characterized. This is partly due to the fact thatit is difficult to positively identify FeS phases using X-raydiffraction (XRD) and microscopic techniques in thesecomplex sediment and soil samples (12). A field study byHerbert et al. (10) demonstrated the inability to identify ironsulfide minerals from reactive media by either XRD orscanning electron microscopy (SEM), and formation of FeSwas identified only on piezometer tubing. In this study, FeSphases were directly identified by XRD and SEM in thereactive material of column and field samples.

In 2002, a biowall composed of sand and plant mulch wasconstructed at Altus Air Force Base, OK to treat a TCE plumein groundwater. A column study was conducted to simulatefield conditions at Altus AFB in order to better understandthe mechanisms responsible for removal of TCE (9). Removalof TCE is extensive (>92-99%) in both the biowall and thecolumn experiments (9, 14). However, the mechanism forTCE degradation in these systems is only partially understood.Based on the extent of sulfate reduction, FeS formation inthe columns could be estimated, and Shen and Wilson (9)concluded that TCE removal associated with FeS couldaccount for about one-half of TCE degradation in the columnexperiment. Acetylene and 13C labeled CO2 constituted themajor products from abiotic reduction of 13C labeled TCE(9). Butler and Hayes (6) showed that TCE can be degradedby FeS through a reductive elimination pathway, producingacetylene as a major reaction product. Acetylene is a signatureof abiotic reduction.

Amorphous or poorly crystalline FeS is more effectivethan FeS2 in degrading TCE (15, 16). To predict the long-term ability of a PRB to remove TCE, it is necessary to havea better understanding of the processes that form FeS andthe processes that transform FeS to FeS2 in PRBs. We usedacid volatile sulfide (AVS) and chromium reducible sulfur(CRS) as probes for FeS formation and transformation toFeS2, and applied XRD and microscopy techniques toelucidate the biogeochemical processes that are relevantto FeS formation and transformation in a PRB.

Materials and MethodsBackground. Figure 1a shows the location of the OU-1 biowallwhich was constructed to treat a TCE groundwater plume atAltus AFB, OK. It is 138.7 m long, 7.3 m deep, and 0.46 mwide. It contains a blend of 50% (v/v) shredded tree mulch,10% (v/v) cotton gin trash, and 40% (v/v) sand. On a volumebasis, the biowall averages 71% solids and 29% water. Theestimated residence time of groundwater in the biowall is 10days.

Shen and Wilson (9) constructed Column B3 (Figure 1b)using the same materials as the OU-1 biowall, with thefollowing exception. While the biowall was constructed with40% sand, Column B3 contained 36% sand and 4% hematite.Hematite was added to Column B3 to determine the benefitof adding a source of reactive Fe(III) in future biowalls. Shenand Wilson (9) had constructed four columns simulating aplant mulch biowall. Based on a mass balance of sulfateentering the columns and sulfate and sulfide leaving thecolumns, Column B3 had accumulated the highest concen-tration of FeS. Column B3 was selected for a detailedcharacterization of the fractionation of iron and sulfur mineralphases.

* Corresponding author phone:580-436-8554; fax: 580-436-8703;e-mail: he.yongtian@epa.gov.

† National Research Council.

Environ. Sci. Technol. 2008, 42, 6690–6696

6690 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 42, NO. 17, 2008 10.1021/es8010354 CCC: $40.75 2008 American Chemical SocietyPublished on Web 08/05/2008

The column was supplied with groundwater from AltusAFB amended with TCE to a final concentration of 15 µM.The water was pumped upward from the inlet at the bottomof the column with a peristaltic pump. The mean residencetime of groundwater in the column was 17 days. The chemicalcomposition of materials used to construct Column B3 andthe field scale OU-1 biowall are listed in Table S1 and S2 inthe Supporting Information (SI).

Groundwater and Solid Phase Sampling. Procedures forsampling and for chemical analysis of water samples aredescribed in the SI. Figure S1 in the SI shows the locationsof monitoring wells at the OU-1 site at Altus AFB.

After 875 days of operation, Column B3 was put into afreezer for 1 week. This was done to immobilize the contentsof the column and protect them from oxidation duringexposure to the atmosphere, and to break the glass containeraway from the frozen column of plant mulch, sand, andgroundwater. The broken glass was discarded and the columnwas cut while still frozen into 5 cm thick sections using ahand saw. Each section was thawed in an anaerobic glovebox(N2:H2 ) 92.5%:7.5%), homogenized and separated into fourequal subsections and refrozen until analysis. Samples thatwere extracted and analyzed to determine iron and sulfurpartitioning were thawed in the glovebox in sealed containers.Wet sediments were used directly for analysis and correctedfor H2O content. H2O content was determined gravimetrically.For total sulfur measurement, samples were dried inside theglovebox.

To acquire samples of the OU-1 biowall, a freeze coresampling technique was developed. This technique usedliquid N2 to freeze mulch material in the biowall in situ. Mulchmaterials were then retrieved as a “Popsicle” (Figure S2 inSI). Each core collected from the subsurface was immediatelywrapped in a plastic bag and placed on dry ice in an icechest. Further sample preparation for iron and sulfur analysisfollowed the same procedure as the column samples.

Iron and Sulfur Partitioning. The iron partitioningprocedure was modified after Kosta and Luther (17). For sulfurpartitioning, the extraction and experimental setup followedprocedures in Wilkin and Bischoff (18). Table 1 provides thedetails of each extraction step and the targeted iron mineralphases.

Solid Phase Characterization. X-Ray diffraction (XRD)scans were collected with a Rigaku MiniFlex diffractometer(Fe KR radiation, operated at 30 keV and 15 mA). As apreconcentration step to prepare XRD samples, mulch

materials were washed with methanol, and large sandparticles and splinters of mulch were removed. A fine-grainedfraction was collected using a 0.4 µm filter and dried in avacuum desiccator.

Scanning electron microscope (SEM) images were ob-tained using a JEOL JSM-6360 scanning electron microscopewith EDS/Oxford INCA software. SEM samples were preparedby placing a drop of methanol suspension of mulch on acarbon planchet (Electron Microscopy Sciences). The planchetwas placed in a vacuum desiccator to dry overnight, andthen coated with gold prior to SEM examination.

Equilibrium geochemical modeling software (Geochem-ist’s Workbench, Release 6.0, Rockware, Golden, CO) wasused to model groundwater chemical speciation. Modelsimulations considered the formation of solution complexesand precipitation of solids. The solubility model of Rickard(19) was used to calculate saturation indices of FeS.

Results and DiscussionGroundwater Chemistry. The operating conditions forColumn B3 simulated groundwater conditions at the OU-1biowall of Altus AFB. Typical chemical compositions ofinfluent and effluent from Column B3 are shown in Table S3in the SI. In general, the column effluent contained lowerconcentrations of dissolved oxygen (DO) and sulfate andhigher concentrations of dissolved sulfide and alkalinitycompared to the influent. Alkalinity was increased due toCO2 production during biodegradation of the plant mulch.In column B3, concentrations of sulfide was never higherthan 16 mg/L. Typical composition of groundwater at theAltus biowall field site is shown in Table S4 in the SI. Insidethe biowall, sulfide concentrations range from 60 to 220 mg/Ldepending on location.

Hematite was not added during construction of the fieldscale OU-1 biowall. There was a limited amount of iron inthe sand added to the mulch material, and as a result thepotential for FeS formation in the OU-1 biowall wassignificantly lower than in Column B3, which explains thehigher concentrations of sulfide measured in groundwaterin the field compared to the effluent of Column B3. The H2Sconcentrations in the system are controlled by bacterialsulfate reduction and the availability of reactive iron.

Geochemical modeling indicates that groundwater withinthe biowall is near saturation with respect to iron mono-sulfide. Mean saturation indices for crystalline mackinawiteand disordered mackinawite are 0.38 and-0.41, respectively

FIGURE 1. (a) Location of a PRB (a biowall) constructed at Altus AFB in June 2002. (b) Column B3 constructed to simulate fieldconditions in the OU-1 biowall.

VOL. 42, NO. 17, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 6691

(n ) 17). These results indicate that geochemical modelingbased on water chemistry parameters is useful in evaluatingwhether reactive mineral phases might be present at sitesunder situations where solids cannot be collected and directlyanalyzed for reactive mineral phases.

Solid Phase Iron Fractionation. Figure 2 presents dataon iron partitioning in Column B3 and in core samples fromthe OU-1 mulch biowall. Iron extracted by ascorbate includesamorphous Fe(III) oxides. The concentration of iron extract-able by ascorbate was negligible. Iron extracted by HClincludes both amorphous Fe(III) oxides and iron associatedwith AVS. The low concentration of iron extracted byascorbate and the high concentration of iron extracted byHCl in material from Column B3 suggests that almost all theiron was reduced under anaerobic conditions in the columnafter 875 days of continuous operation. The low concentrationof iron extacted by HCl in the biowall samples suggests thatthe original pool of available reactive iron was small comparedwith Column B3. Dithionite extracts amorphous Fe(III)

oxides, crystalline Fe(III) oxides and iron associated withAVS. Measured concentrations of iron extracted with HCl orwith dithionite are very similar across the column and OU-1field samples, which indicates that there are only traceamounts of crystalline Fe(III) oxides remaining in thesematerials. Total iron was measured as iron extracted by HNO3

during microwave digestion. In the biowall samples and thefew sections of Column B3 that are further down gradientof the inlet, the concentration of total iron is significantlyhigher than iron that was extacted by HCl or dithionite. Totaliron includes the FeS2 fraction, which is not extracted by HClor dithionite.

Reactive iron has traditionally been defined as that fractionof iron in sediments which reacts readily with sulfide to formiron sulfides and pyrite (17, 20). The iron which may reactwith H2S (reactive iron) is mainly iron that is not bound insilicate minerals, including amorphous iron (oxyhydr)oxides,some crystalline iron oxides, and iron monosulfides. Op-erationally, reactive iron has been defined as iron extracted

TABLE 1. Extraction Methods Used to Characterize the Fractionation of Iron and Sulfur Species in the Samples

extraction extractants pH time target mineral phases

iron extraction proceduremodified after Kostkaand Luther (17).

ascorbate 10 mL extractant composed of 0.17 Msodium citrate (Aldrich),0.60 M sodium bicarbonate (Fisher) and0.11 M ascorbic acid (Fisher)

8 24 h amorphous iron oxides

HCl 10 mL 0.5 M HCl (J.T.Baker) <2.0 1 h amorphous iron oxides, AVSdithionite 10 mL extractant composed of 0.29 M

sodium dithionite (Fisher),0.35 M sodium acetate (J.T.Baker) and0.2 M sodium citrate at 60 °C

4 4 h amorphous iron oxides, crystallineiron oxides, AVS

HNO3 Microwave digestion at 175 °C in 10%HNO3 (JT Baker)

0.5 h total iron

Sulfur fractionationprocedure after Wilkinand Bischoff (18)

AVS 1 M HCl (J.T. Baker) amorphous FeS, mackinawite,greigite

CRS 1 M CrCl2 in 0.5 M HCl,Prepared by passing 1 MCrCl3.6H2O (Fluka, 98+%) dissolvedin 0.5 M HCl through a standardJones redactor

elemental S, pyrite

Total sulfur V2O5 (Aldrich,98+%), sample combustedin two furnaces aligned in sequence(UIC model CM3220)

AVS, CRS, SO42-

FIGURE 2. Iron fractionation in sections of Column B3 and core samples from the OU-1 biowall. (a) column B3, (b) OU-1 fieldsamples.

6692 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 42, NO. 17, 2008

by HCl (21). In the OU-1 biowall, the pool of remainingreactive iron is small compared to the pool of total iron,between 480 and 1000 mg/kg of reactive iron compared to1270 to 2780 mg/kg total iron. In Column B3, the pool ofreactive iron is large compared to the biowall, extending from1700 to 4930 mg/kg in Column B3, compared to 480 and1000 mg/kg in the biowall.

Solid Phase Sulfur Fractionation. AVS can form duringreductive dissolution of iron hydroxides or by reaction offerric iron in clay minerals during reduction with sulfide (22).AVS is operationally defined as reduced inorganic sulfur thatreacts with HCl to form H2S gas (23). The AVS procedureemployed in this study was intended to digest labile sulfidecompounds; it recovers sulfide in pore water and solid-phasesulfide minerals, which include amorphous FeS, mackina-wite (FeS0.9-1.1) and greigite FeS1.33 (Fe3S4) (23, 24).

The CRS fraction has been defined as sulfide phases thatare not extractable in HCl over short time-frames (25). TheCRS extraction procedure (by CrCl2-HCl) is intended to digestrefractory sulfur-bearing compounds, such as pyrite/mar-carsite (FeS2) and elemental sulfur (S0).

Figure 3a shows the relative distribution of AVS, CRS, totalreduced sulfur (TRS, TRS ) AVS + CRS) and total sulfur ateach section of Column B3. CRS is high in the influent section(>2000 mg/kg). There is little CRS in the 10-15 cm segment,and CRS increases in upper sections (20-46 cm). In thesesections, AVS and CRS seem to be inversely correlated.

Figure 3b shows sulfur partitioning across depths of thebiowall at the Altus field site. Compared to Column B3,measured AVS and CRS concentrations in the Altus biowallare much lower.

The low concentration of solid phase sulfide suggests thatthe low concentrations of reactive iron in the sediment limitedFeS formation. In shallow depths of 1.52-1.98 m and1.98-2.44 m, the CRS/AVS ratio is higher than samples fromdeeper depths (>3.51 m). At deeper depths, less AVS istransformed to CRS. Even though the sulfide concentrationinside the wall is high, it apparently does not precipitate aselemental S or form other intermediate sulfur species due tolack of available oxidants (presumably dissolved O2) at deeperdepths. These intermediate species are believed to beessential in the transformation of AVS (FeS) to CRS (FeS2).

Sulfur partitioning in column and field samples show thatAVS and CRS coexist, but the ratio of AVS/CRS is related tolocal and micro geochemical conditions. Over time, sedi-mentary AVS tends to be replaced by the thermodynamicallymore favorable CRS fraction (pyrite) (23). Rickard and others(26) have shown that the kinetics of iron monosulfideprecipitation are fairly rapid, whereas the transformation ofa FeS precursor phase to pyrite is a slower process.

Concentrations of AVS and CRS in Column B3 are inverselyrelated. This suggests that at least part of the CRS is producedfrom AVS. Previous studies suggest that the conversion ofAVS to CRS occurs mostly at redox boundaries where anoxidant (e.g., O2) is available (27). This is consistent with ourobservations in the column experiments, that transformationof AVS to CRS was more extensive in the influent end of thecolumn where potential oxidants such as dissolved oxygenwere available.

Based on the sulfur partitioning data (Figure 3), and theassumption that all the CRS in the samples was originallyAVS, it is possible to calculate the extent of transformationof AVS to CRS in the samples. Data are presented in TablesS5 and S6 in the SI. In the biowall samples, from 55 to 58%of the sulfide remains as AVS. In the sections of Column B3,from 87 to 19% of the sulfide remains as AVS, with an averageof 45%. Within 5.2 years for the biowall, and 2.4 years forColumn B3, roughly one-half of the reactive FeS wastransformed to the much less reactive FeS2. The estimatedrate of transformation of AVS to CRS in the influent sectionof Column B3 is 1.5 × 10-3 mol kg-1 d-1, which is similar toa transformation rate reported by Canfield et al. (28). Theysuggested bacteria in these systems could play a role in theprocess of conversion of FeS (AVS) to FeS2 (CRS).

Mineralogical Characteristics. Mineralogy of the biowallmaterial in column experiments and field core samples wascharacterized by XRD and SEM/EDX. In the coarse sedimentsamples, quartz dominates the sediment mineralogy. Inmethanol washed samples, in addition to quartz, pyrite wasidentified by both XRD and SEM, and residual hematite wasobserved in Column B3 samples (Figure 4, and SI Figure S3and S4).

Under anaerobic conditions, formation of amorphous ironsulfide and mackinawite could occur as finely dispersedprecipitates and as coatings on hematite and other minerals(23). FeS formation was not confirmed by XRD, but isidentifiable by EDX based on chemical composition inColumn B3 samples. The EDX data show that the metastableiron sulfides have a chemical composition of FeS0.9-1.4, andprobably represent a mixture of stoichiometric mackinawiteand greigite. The FeS/greigite morphology (Figure 4b) issimilar to chemically synthesized FeS (unpublished data, Heet al.). This noncrystalline FeS is generally the first phase ofFeS to precipitate. In heterogeneous mixtures, it is difficultto positively identify FeS minerals with XRD due to lowcrystallinity of the FeS phases. Contrary to the experiencewith the samples from Column B3, FeS or greigite phasescould not be identified in the field sediments by either XRDor SEM/EDX.

FIGURE 3. Sulfur fractionation in sections of Column B3 and core samples from the OU-1 biowall. (a) Column B3, (b) OU-1 biowall.

VOL. 42, NO. 17, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 6693

Pyrite was the only FeS phase confirmed by XRD, andEDX analysis of the reaction products returned Fe:S ratios(1:2) consistent with FeS2 stoichiometry. SEM showeddifferent morphologies of pyrite framboids (Figure 4 andFigure S5 in the SI). Some framboids show unorderedmicrocrystal arrangement, other framboids show differentshapes of microcrystals (triangular, round and rectangular),and framboids are present in random sizes. Framboiddiameter size ranges from about 6 to 20 µm. Butler andRickard (29) suggest that differences in framboid morphologycould be related to formation kinetics. In the OU-1 fieldsamples, SEM also shows organic coatings formed onframboid surfaces. This could be an indication that bacteriaplay a role in pyrite framboid formation.

Iron Sulfide Formation and Transformation.FeS Formation. In the OU-1 biowall and in the column

experiments, sulfate reduction is mediated by sulfate reducingbacteria (SRB) under anaerobic conditions. Accompanyingsulfate reduction in the system, reductive dissolution ofhematite and other iron oxides can occur through twopathways: dissimilatory iron reduction by iron reducingbacteria and reductive dissolution of hematite by hydrogensulfide (22).

In the column experiment and the OU-1 biowall, mea-sured sulfide concentrations range from 6 to 228 mg/L. Theproduction of dissolved sulfide species in the presence of

dissolved metals can result in precipitation of metal sulfidemineral phases such as FeS.

H2S+ Fe2+f FeS+ 2H+ (1)

Other metals such as Mn, Ni, and Zn, can also be releasedto groundwater under iron and sulfur reducing conditions,and can react with sulfide to form sparingly soluble sulfideminerals. SEM/EDX showed NiS and ZnS formation in thesystem, but they are only minor components compared toFeS because there were only trace amounts of these elementspotentially available to react with sulfide (Table S2 in the SI).

Transformation of FeS to FeS2. FeS is redox sensitive andmetastable in the environment. It may be transformed toFeS2 based on several different mechanisms. A polysulfidepathway is proposed to be the primary pathway for formationof FeS2(s) in suboxic sediments. In this process iron sulfideprecursors react with S0, polysulfides or other sulfur inter-mediates to form pyrite. Production of intermediate sulfurspecies (for example, elemental sulfur, polysulfides andthiosulfates (S2O3

2-)) is initiated through H2S oxidation nearthe redox boundary by oxidants such as O2, NO3

-, and ironoxides and manganese oxides (29, 30). Transformation viareaction with elemental sulfur is most likely the dominantpathway in our systems based on the observation that greater

FIGURE 4. SEM of Column B3 samples. (a). Hematite, FeS, and FeS2 particles; b) FeS, (c) pyrite framboids. d, e, and f arebackscattered electron images of the OU-1 biowall samples. (d) framboids showing no clear microcrystals; (e) framboid anddispersed microcrystals; (f) pyrite framboid with surface coating.

6694 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 42, NO. 17, 2008

conversion to CRS occurred only in sections of Column B3where potential oxidants are available.

Under strictly anoxic conditions, FeS2(s) has also beenshown to form via a second mechanism, the “H2S oxidationpathway” (26, 31).

FeS(s)+H2S(aq)f FeS2(s)+H2(g) (2)

Rickard and Luther (31) show the H2S pathway to proceedmore rapidly than the polysulfide pathway and all otherpathways involving intermediate sulfur species. High sulfideconcentration in the field and column is conducive to thismechanism.

A third mechanism, the “direct” pyrite precipitationpathway, could occur through prolonged exposure of iron-containing minerals to dissolved sulfide. Pyrite originatesfrom exposure of iron(II)-containing minerals to dissolvedsulfide and most likely precipitates on a pyrite precursor(“seed”) without the formation of a FeSn precursor (30). Thedirect pyrite formation from iron oxides may be describedby the following reaction (18):

2Fe(OH)3 + 3HS-f FeS+ FeS2+3H2O+ 3OH- (3)

For the first pathway, intermediate sulfur species arecritical in pyrite formation. Thus the availability of an oxidantto produce sulfur species with intermediate oxidation statescontrols the transformation of FeS to pyrite. The transfor-mation of FeS to FeS2 is most pronounced at the redoxinterface in our systems. The conversion of iron monosulfidesto pyrite in Column B3 occurs in the section adjacent to theinlet as well as the down gradient parts of the columns, andan increase in CRS is accompanied by a decrease in AVS.Higher levels of CRS in the section adjacent to the inlet couldbe due to dissolved O2 in the influent. Water analyses inColumn B3 showed that dissolved O2 level is greatly reducedin the effluent, presumably consumed in the bottom sectionof Column B3. The increased CRS concentration in the middleand far down gradient sections of the column could be dueto high concentrations of S0 in the down gradient sections,which facilitated pyrite formation. In a PRB field study,Herbert (10) observed a greater accumulation of mildlyoxidized sulfur species (pyrite, S0) at the front end of thebarrier. This observation suggests that oxidizing agents maybe present in the upgradient aquifer in low concentrations,but are consumed as groundwater passes through the reac-tive barrier.

Implications. The formation of reactive FeS is importantto the degradation of chlorinated solvents in permeablereactive barriers. Shen and Wilson (9) show that abiotictransformation of TCE by FeS could provide a majorcontribution to the removal of TCE from the groundwaterat the Altus AFB. Abiotic degradation of TCE associated withFeS may have an important role in many natural attenuationand engineered remedies.

FeS is metastable and will transform to more stable FeS2

over time. FeS to FeS2 transformation in biowalls couldpotentially reduce the rate of TCE degradation. Weerasooriyaand Dharmasena (15) show FeS2 is much less efficient inTCE degradation (15-66 times slower). Lee and Batchelor(16) indicated that pyrite degraded TCE at a rate>1000 timesslower than FeS. Thus the pyritization process observed inthe column and field experiments in this study can signifi-cantly decrease the extent of degradation of TCE in PRBs.

A supply of reactive iron in the matrix of the biowall willstimulate the formation of FeS. This FeS will promote theabiotic dechlorination pathway (5, 6), as opposed to biologicalreductive dechlorination. In the abiotic degradation pathway,much of the TCE is degraded to acetylene, which avoids anypotential for the formation of vinyl chloride, which is a majorintermediate in biological degradation.

Engineered systems that are designed to exploit abioticdegradation will be most successful when the concentrationsof naturally occurring reactive iron are high or whensupplemental reactive iron is introduced to the subsurface.

AcknowledgmentsThe U.S. Environmental Protection Agency through its Officeof Research and Development funded the research describedhere. It has not been subjected to Agency review and thereforedoes not necessarily reflect the views of the Agency, and noofficial endorsement should be inferred. Y.T. He is anassociate of the National Research Council.

Supporting Information AvailableAddtional information on sampling and analysis of ground-water, location of monitoring wells, XRD data, SEM micro-graphs, chemical composition of materials used for con-struction of the biowall and the laboratory column,geochemical parameters in water from the column and thebiowall, and the iron and sulfur fractionation in samplesfrom the column and biowall.This material is available freeof charge via the Internet at http://pubs.acs.org.

Literature Cited(1) Arnold, W. A.; Roberts, A. L. Pathways and kinetics of chlorinated

ethylene and chlorinated acetylene reaction with Fe(0) particles.Environ. Sci. Technol. 2000, 34, 1794–1805.

(2) Amonette, J. E.; Workman, D. J.; Kennedy, D. W.; Fruchter, J. S.;Gorby, Y. A. Dechlorination of carbon tetrachloride by Fe(II)associated with goethite. Environ. Sci. Technol. 2000, 34, 4606–4613.

(3) Lee, W.; Batchelor, B. Abiotic reductive dechlorination ofchlorinated ethylenes by iron-bearing soil minerals. 2. Greenrust. Environ. Sci. Technol. 2002, 36, 5348–5354.

(4) Danielsen, K. M.; Hayes, K. F. pH dependence of carbontetrachloride reductive dechlorination by magnetite. Environ.Sci. Technol. 2004, 38, 4745–4752.

(5) Butler, E. C.; Hayes, K. F. Kinetics of the transformation oftrichloroethylene and tetrachloroethylene by iron sulfide.Environ. Sci. Technol. 1998, 32, 1276–1284.

(6) Butler, E. C.; Hayes, K. F. Effects of solution composition andpH on the reductive dechlorination of hexachloroethane byiron sulfide. Environ. Sci. Technol. 1999, 33, 2021–2027.

(7) Kennedy, L. G.; Everett, J. W.; Gonzales, J. Assessment ofbiogeochemical natural attenuation and treatment of chlori-nated solvents, Altus AFB, Altus, Oklahoma. J. Contam. Hydrol.2006, 83, 221–236.

(8) Kennedy, L. G.; Everett, J. W.; Becvar, E; DeFeo, D. Field-scaledemonstration of induced biogeochemical reductive dechlo-rination at Dover AFB, Dover, Delaware. J. Contam. Hydrol.2006, 88, 119–136.

(9) Shen, H.; Wilson, J. T. Trichloroethylene removal from ground-water in flow-through columns simulating a permeable reactivebarrier constructed with plant mulch. Environ. Sci. Technol.2007, 41, 4077–4083.

(10) Herbert, R. B.; Benner, S. G.; Blowes, D. W. Solid phase iron-sulfur geochemistry of a reactive barrier for treatment of minedrainage. Appl. Geochem. 2000, 15, 1331–1343.

(11) Roh, Y.; Lee, S. Y.; Elless, M. P. Chracterization of corrosionproducts in the permeable reactive barriers. Environ. Geol. 2000,40, 184–194.

(12) Phillips, D. H.; Gu, B. W. D. B.; Roh, Y.; Liang, L; Lee, S. Y.Performance evaluation of a zerovalent iron reactive barrier:Mineralogical characteristics. Environ. Sci. Technol. 2000, 34,4169–4176.

(13) Benner, S. G.; Blowes, D. W.; Gould, W. D.; Herbert, D. H.; Ptacek,C. J. Geochemistry of a permeable reactive barrier for metalsand acid mine drainage. Environ. Sci. Technol. 1999, 33, 2793–2799.

(14) Lu, X.; Wilson, J. T.; Shen, H.; Henry, B. M.; Kampbell, D. H.Remediation of TCE-contaminated groundwater by a permeablereactive barrier filled with plant mulch (Biowall). J. Environ.Sci. Health A 2008, 43, 24–35.

(15) Weerasooria, R.; Dharmasena, B. Pyrite-assisted degradationof trichloroethene (TCE). Chemosphere 2001, 389–396.

VOL. 42, NO. 17, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 9 6695

(16) Lee, W.; Batchelor, B. Abiotic reductive dechlorination ofchlorinated ethylenes by iron-bearing soil minerals. 1. Pyriteand magnetite. Environ. Sci. Technol. 2002, 36, 5147–5154.

(17) Kostka, J. E.; Luther, G. W. Partitioning and speciation of solidphase iron in saltmarsh sediments. Geochim. Cosmochim. Acta1994, 58, 1701–1710.

(18) Wilkin, R. T.; Bischoff, K. J. Coulometric determination of totalsulfur and reduced inorganic sulfur fractions in environmentalsamples. Talanta 2006, 70, 766–773.

(19) Rickard, D. The solubility of FeS. Geochim. Cosmochim. Acta2006, 70, 5779–5789.

(20) Berner, R. A. Sedimentary pyrite formation. Am. J. Sci. 1970,268, 1–23.

(21) Burton, E. D.; Phillips, D. Reactive sulfide relationships withtrace metal extractability in sediments from southern MoretonBay. Aus. Mar. Pollut. Bull. 2005, 50, 589–608.

(22) Pyzik, A. J.; Sommer, S. E. Sedimentary iron monosulfides:Kinetics and mechanisms of formation. Geochim. Cosmochim.Acta 1981, 45, 687–698.

(23) Morse, J. W.; Cornwell, J. C. Analysis and distribution of ironsulfide minerals in recent anoxic sediments. Mar. Chem. 1987,22, 55–69.

(24) Morse, J. W.; Rickard, D. Chemical dynamics of sedimentaryacid volatile sulfide. Environ. Sci. Technol. 2004, 38, 131A–136A.

(25) Huerta-Diaz, M. A.; Morse, J. W. A quantitative method fordetermining trace metal concentrations in sedimentary pyrite.Mar. Chem. 1990, 29, 119–144.

(26) Rickard, D. Kinetics of pyrite formation by the H2S oxidationof Fe(II) monosulfide in aqueous solutions between 25 and125�C: the rate equation. Geochim. Cosmochim. Acta 1997, 61,115–134.

(27) Wilkin, R. T.; Barnes, H. L. Pyrite formation by reactions of ironmonosulfides with dissolved inorganic and organic sulfurspecies. Geochim. Cosmochim. Acta 1996, 60, 4167–4179.

(28) Canfield, D. E.; Thamdrup, B.; Fleischer, S. Isotopic fractionationand sulfur metabolism by pure and enrichment cultures ofelemental sulfur-disproportionating bacteria. Limnol. Oceanogr.1998, 43, 253–264.

(29) Butler, I. B.; Rickard, D. Framboidal pyrite formation via theoxidation of iron(II) monosulfide by hydrogen sulfide. Geochim.Cosmochim. Acta 2000, 64, 2665–2672.

(30) Schoonen, M. A. A.; Barnes, H. L. Reactions forming pyrite andmarcasite from solution: II. via FeS precursors below 100° C.Geochim. Cosmochim. Acta 1991, 55, 1505–1514.

(31) Rickard, D.; Luther, G. W. Kinetics of pyrite formation by theH2S oxidation of iron(II) monosulfide in aqueous solutionsbetween 25 and 125° C: the mechanism. Geochim. Cosmochim.Acta 1997, 61, 135–147.

ES8010354

6696 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 42, NO. 17, 2008